Calculate the emf and ΔG of the cell reaction for the following cell at 25°C: $$\mathrm{Mg(s)\,|\,Mg^{2+}(0.001\ M)\,||\,Cu^{2+}(0.0001\ M)\,|\,Cu(s)}$$ ($E^{\circ}_{\mathrm{Mg^{2+}/Mg}} = -2.37$ V; $E^{\circ}_{\mathrm{Cu^{2+}/Cu}} = +0.34$ V; 1 F = 96500 C mol$^{-1}$)

CBSE Class 12 · 2000 · Delhi · Redox Reactions and ElectrochemistryShort Answer

Calculate the emf and ΔG of the cell reaction for the following cell at 25°C:

Mg (s) ∣ M g^{2 +} (0.001 M) ∣∣ C u^{2 +} (0.0001 M) ∣ Cu (s)

(

E_{M g^{2 +} /Mg}^{\circ} = - 2.37

E_{C u^{2 +} /Cu}^{\circ} = + 0.34

V; 1 F = 96500 C mol

^{- 1}

)

Answer: E_cell ≈ 2.68 V; ΔG ≈ −517 kJ mol⁻¹

E_{ce l l}^{\circ} = 0.34 - (- 2.37) = 2.71 V, n = 2

Nernst equation:

E_{ce l l} = 2.71 - \frac{0.059}{2} lo g \frac{[ M g ^{2 +} ]}{[ C u ^{2 +} ]} = 2.71 - 0.0295 lo g \frac{1 0 ^{- 3}}{1 0 ^{- 4}}

= 2.71 - 0.0295

E_{ce l l} = 2.68 V

Gibbs energy:

Δ G = - n F E_{ce l l} = - 2 \times 96500 \times 2.68

Δ G = - 5.17 \times 1 0^{5} J = - 517.2 kJ mo l^{- 1}

CBSE Class 12 Board · 2000 · Delhi · Subjective